Thermodynamic Modeling and Experimental Study of CO2 Dissolution

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Thermodynamic Modeling and Experimental Study of CO2 Dissolution in New Absorbents for Post-Combustion CO2 Capture Processes Yohann Coulier, Alexander R Lowe, Jean-Yves Coxam, and Karine Ballerat-Busserolles ACS Sustainable Chem. Eng., Just Accepted Manuscript • DOI: 10.1021/ acssuschemeng.7b03280 • Publication Date (Web): 30 Nov 2017 Downloaded from http://pubs.acs.org on December 11, 2017

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ACS Sustainable Chemistry & Engineering

Thermodynamic Modeling and Experimental Study of CO2 Dissolution in New Absorbents for Post-Combustion CO2 Capture Processes Y. Couliera*, A. R. Lowea, J-Y. Coxama, K. Ballerat-Busserollesa,b a)

Université Clermont Auvergne, CNRS, SIGMA Clermont, Institut de Chimie (ICCF), F63000 Clermont–Ferrand, France. b) Mines ParisTech PSL – Research University, CTP, 35 rue Saint Honoré, 77305, Fontainebleau * Corresponding author: [email protected]

Abstract Demixing solvents have been suggested as an alternative to aqueous solutions of alkanolamines for the decarbonation of industrial effluents. These solvents undergo a liquidliquid phase separation when heated. A phase separation step could be integrated in carbon dioxide (CO2) capture process to reduce energetic cost of the solvent regeneration step. As part of a general project on the dissolution of CO2 in aqueous solutions of piperidines, 3methylpiperidine and 4-methylpiperidine has been investigated. The enthalpies of solution of CO2 in aqueous solutions of those methylpiperidines have been measured using a flow calorimetric technique. This technique makes it possible to simultaneously determine the enthalpies of solution and CO2 solubility limits. The experimental data were used to develop a thermodynamic model representative of phase equilibrium and enthalpy of solution. The carbamate formation constants of the 3-methylpiperidine and 4-methylpiperidine are adjusted in the model. The thermodynamic model is used to discuss the mechanism of CO2 dissolution.

Keywords Phase change solvent; 3-methylpiperidine; 4-methylpiperidine; enthalpy of solution; solubility limit; equilibrium constant for carbamate formation

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INTRODUCTION Removing carbon dioxide (CO2) from post-combustion industrial effluents will contribute substantially to the reduction of anthropogenic emission of carbon. Effective CO2 emission mitigation strategies such as Carbon Capture and Storage (CCS) are required to tackle climate change.1 Amine based solvents, such as aqueous solutions of alkanolamines, have been used for the last six decades in oil and gas industries to separate CO2 from gas streams. The technology could be applied to CO2 removal from post-combustion effluents in power plants.2 Despite new challenges due to the unique composition of flue gasses and low CO2 partial pressures, the current energy demand to regenerate the amine solution remains a main economic barrier.3,

4

Energy efficiency is thus considered as a key challenge for the

development of carbon capture technologies for greenhouse gas reduction5,

6

and research

efforts are focused on the identification of breakthrough processes that will reduce the energy consumption.7 Demixing solvents (DMX)8 have been proposed as an alternative to classical alkanolamines. In new processes such as the DMX™ process developed by IFP Energies Nouvelles,9 a decantation step is added before the CO2 desorption step. The liquid-liquid phase separation make it possible to treat only the CO2-rich phase in the desorption unit; the CO2-lean phase is re-injected into the absorption unit. The new absorbents must be able to demix according to particular conditions of temperatures (higher than the absorption temperature) and CO2 loading charges. This phase separation aims to concentrate CO2 into the water rich phase, lowering thus the amount of absorbent to be treated in the regeneration step and reducing the energy requirements. The DMX™ technology was demonstrated in an industrial pilot project10 and a lowering around 22 % of the general energy requirements was reported. Other processes using liquid-liquid phase change solvent properties have also been developed during the last decade.11 For example, a selective CO2 absorption with phase separation can also be obtained with thermomorphic biphasic solvent (TBS)12 or mixed phase-changed solvents.13 The {CO2 – DMX solvent} system can be studied using a thermodynamic approach. The purpose is to develop models adapted to correctly estimate the absorption properties required to design and dimension process units. Due to the complexity of the studied system, thermodynamic modeling needs numerous and various experimental data. Liquid-vapor and liquid-liquid equilibrium data were obtained for different DMX solvents (DMX absorbents are not disclosed) by Aleixo et al..14 Coulier et al.15-17 found that N-methylpiperidine and 2methylpiperidine (2MPD) show a partial miscibility with water as observed with the DMXTM absorbents. Then a general study on substituted piperidine molecules has been carried out to ACS Paragon Plus Environment

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investigate the influence of the addition of alkyl groups on the CO2 absorption and phase separation properties. The present work reports the results obtained with two methylpiperidine isomers: 3-methylpiperidine (3MPD) and 4-methylpiperidine (4MPD). It has been initially shown18-20 that the temperature of the liquid-liquid equilibrium depends on the position of the methyl group on the piperidine ring. Indeed the lower critical solution temperatures (LCST) were found to be 339 K at a molar fraction (xa) of 0.04 for 3MPD and, 360 K and xa = 0.06 for 4MPD.18-20 The aim of the current research is to provide experimental thermodynamic data on carbon dioxide dissolution in aqueous solution of 3-methylpiperidine and 4-methylpiperidine. This paper will focus on the heat of gas dissolution. This quantity is of interest to estimate the energy required for regenerating the CO2 loaded amine solution and is thus a key parameter for the design of CO2-removal unit. The mechanism of CO2 capture by absorption in aqueous amine solutions is governed by a combination of physical dissolution and chemical reactions between CO2 and aqueous amine solution. The enthalpy of solution measured when mixing CO2 and aqueous solution of demixing amine is the total heat effects due to both the physical dissolution and the chemical reactions. To our knowledge, there is no current literature reporting energetic properties of CO2 reacting with these two cyclic amines. The enthalpies of solution of CO2 in both amine based systems were determined for mass fraction of amine ranges from wa = 0.17 to 0.4, at temperatures of 318 K and 338 K and pressures from 0.5 MPa to 1.5 MPa. Experimental conditions have been chosen to avoid phase separation during gas absorption. A thermodynamic model previously described by Coulier et al.21 was applied to represent the energetic effects during CO2 dissolution. A comparative study of methylpiperidine isomers, 2-methylpiperidine, 3-methylpiperidine and 4-methylpiperidine, based on modeling results is given in this paper.

EXPERIMENTAL SECTION

Chemicals and Materials. 3-methylpiperidine and 4-methylpiperidine were used without further purification. Water was distilled and degassed under vacuum before use (resistivity 18.2 MΩ·cm). Aqueous solutions were prepared by mass; uncertainty in mass fraction (w) is estimated to be less than ± 10-4. Aqueous solutions were stored in glass bottle in an opaque cabinet to prevent any photo-



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degradation. Suppliers, purities and CAS numbers of all chemicals used in this study are given in supporting information in Table S1.

Experimental device and methods. The experimental technique used in this work to measure the enthalpy of solution of CO2 in the absorbent has already been described in details by Arcis et al..22 The enthalpy of solution (∆solH) of CO2 in aqueous solution of amine is determined as a function of gas loadings α (mol CO2 / mol amine), using a flow-mixing cell custom-made for a BT2.15 calorimeter from Setaram. Experiments were performed at constant temperature (± 0.01 K) and pressure (± 0.025 MPa). The absorbent solution and the carbon dioxide mixed into the mixing cell located in the calorimetric block where the heat flux due to gas dissolution is measured. The two fluids flow at constant volume flow rate using two high-pressure syringe pumps (Isco Model 100 DM). In order to maintain constant the molar flow rate, the pumps are temperature controlled at 303.15 ± 0.1 K. The molar flow rates are derived from the volume flow rates of the pumps using fluid densities. The densities of CO2 were calculated using the equation of state from Span and Wagner23 and the densities of aqueous amine solutions are issued from Lowe.18 The enthalpy of solution ∆solH is derived from the calorimetric signal difference, detected by the thermopile surrounding the mixing cell as follows: ∆ sol H / kJ ⋅ mol -1 =

(S M - S BL )

(1)

K ⋅ n& i

where subscript i is for CO2 or amine, and K is a calibration constant used to convert the calorimetric signal (mV) into heat flux (mW). SM represents the thermopile signal during the mixing process and SBL the baseline signal recorded only when the aqueous phase is flowing through the calorimeter.The molar enthalpy is calculated using the molar flow rates n& i (mol.s1

) of the gas (∆solH/kJ.mol-1 of CO2, Figure 1a) or the molar flow rate of amine (∆solH/kJ.mol-1

of amine, Figure 1b). The relative uncertainty on loading charge and enthalpies of solution are less than 1.5% and 5 % respectively.22

RESULTS AND DISCUSSION

Experimental enthalpies of solution. The enthalpy of solution of CO2 in the aqueous amine solutions (w3MPD = 0.19 and 0.40 and, w4MPD = 0.17 and 0.40) was investigated at 318 K and 338 K at constant pressures ranging


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from 0.5 MPa to 1.5 MPa. However, the measurements of the enthalpy of solution of CO2 in the 3MPD solution, w3MPD = 0.19, could not be carried out at 338 K. This temperature was too close to the temperature of phase separation of the amine solution. Experiments were conducted for different gas loading charge (α) up to the saturation of the absorbent solution. The experimental results together with their experimental uncertainties are listed in the supporting information in Tables S2-8. Large exothermic effects were observed for both systems at both investigated temperatures. As an example, the enthalpies of solution versus loading charge determined for the 3MPD (w3MPD = 0.1936) and the 4MPD (w4MPD = 0.17) at T = 318 K are illustrated in Figure 1. When the enthalpy of solution ∆solH is expressed in kJ·mol−1 of CO2 (Figure 1a), the graph exhibits plateaus for the lowest loading charges (α ≤ 0.4), and then the exothermic effect decreases as the loading increases. This typical experimental trend has been previously reported in the literature.21, av

22, 24, 25 −1

enthalpy of solution at low loading charge ∆solH expressed in kJ·mol

The values of

of CO2 are average

values obtained on the plateau, for α ≤ 0.4. They are reported for all investigated experimental conditions in Table 1. For both studied systems, under same conditions of pressure and temperature, a small decrease of ∆solHav is observed as the composition rises. A small diminution of the exothermic effect is also noticed when the temperature or the pressure increases for CO2 dissolution in solutions with same amine composition. The results show that the position of the methyl group has no apparent effect on the ∆solHav. In Figure 1b, experimental enthalpies of solution expressed in kJ·mol−1 of 3MD and 4MPD are plotted as function of loading charge. These experimental enthalpies exhibit three different domains. In the first domain (0 < α ≤ 0.5), ∆solH increases linearly with the loading charge. The value of the slope corresponds to ∆solHav obtained previously (Figure 1a). In the second domain (0.5 ≤ α < s), ∆solH keep rising but with a different slope to reach a plateau. The slope change is due to a modification of the absorption mechanisms for α > 0.5. This change is described by the thermodynamic modeling in the next section. The last domain, where the enthalpy of solution stays constant, is characteristic of solution saturation. The intersection between unsaturated (enthalpy increase) and saturated (plateau) domains yields the solubility of the gas (s). The experimental solubilities in aqueous 3MPD and 4MPD solutions were then graphically determined at 318 K and 338 K for each studied pressures. Values are listed in Table 2 together with their uncertainties. The solubilities of CO2 in the aqueous solutions of 3MPD and 4MPD follow the same trends: a decrease with amine composition and temperature. The first effect (amine composition) can be explained by the



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diminution of water molecules available to solvate CO2 when the amine composition (wa) is increased, while the second effect (temperature) is related to the change in amine basicity with temperature. Under the same conditions of temperature and amine composition, the solubilities (s) are nearly constant for aqueous solutions of 3MD and 4MPD.

Thermodynamic and equilibrium model. The thermodynamic representation of CO2 absorption in aqueous solution 3MPD and 4MPD is based on a model of phase equilibria using a γ-φ approach.21, 26 The chemical and physical equilibria that take place during the CO2 dissolution in aqueous solutions of secondary amines such as 3MPD and 4MPD are expressed by the following reactions: H 2 O ⇌ H + + OH -

(2)

CO 2 + H 2 O ⇌ HCO 3− + H +

(3)

HCO 3− ⇌ CO 32− + H +

(4)

+ R R ′NH +2 ⇌ R R ′NH + H

(5)

RR ′NH + HCO3− ⇌ R R ′NCOO - + H 2 O

(6)

RR ′NH (l) ⇌ RR ′NH (v)

(7)

H 2 O (l) ⇌ H 2 O (v)

(8)

⇌ CO (v) CO (l) 2 2

(9)

Each chemical equilibrium reaction (eq. 2-6) is represented by the following equation:

K = ∏ aiυi = ∏ (mi γ i ) i υ

i

(10)

i

where ai, γi, mi and υi are, respectively, the activity, the activity coefficient, the molality and stoichiometric coefficient of species i. The equilibrium constants for reactions 3 to 6 are represented as function of temperature by the following function: ln K = A (T / K ) + B + C ⋅ ln (T / K )

(11)

Coefficients A, B and C are summarized in table 3 together with literature sources. However, no experimental data were reported in the open literature for the carbamate formation (eq 6) with both amines (3MPD and 4MPD). The A, B and C parameters in eq 11 for carbamate formation constant were thus treated as adjustable parameters in the model. Details on the thermodynamic model used in this study are given in the Section B of the Supporting Information. Briefly, the Soave-Redlich-Kwong equation of state27 is used for vapor phase


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non-idealities of the molecular species involved in the vapor-liquid equilibria (eq. 7-9). The activity coefficients of the different species in the liquid are estimated using a modified Pitzer equation28 with the interaction parameters determined by Coulier et al..21 The thermodynamic model was adjusted using the experimental CO2 solubility and enthalpy data determined in this work. The results show a good overall correlation. The agreements between calculated CO2 solubilities in aqueous solutions of 3MPD and 4MPD with the experimental values (s) are depicted in figure 2a and figure 2b, respectively. The model aptitude to assess the enthalpy of solution (∆solH) as function of the gas loading charge is represented in figures 3 and 4. The enthalpy of solution is calculated in our model as a sum of enthalpy contribution terms24, 26 representative of the reaction mechanism steps. Thus the plots in figures 3 and 4 represent the total enthalpy of solution together with enthalpy contribution due to amine protonation, carbamate formation, CO2 physical dissolution, first and second ionizations of CO2. For 3MPD solutions (figure 3), the agreement between calculated enthalpies and experimental results is within experimental uncertainty (5%); the calculated enthalpy of solution is slightly underestimated at 318 K at loading charges α > 0.5, with a maximum deviation of 10% for dissolution in the solution of composition w3MPD = 0.1936. For 4MPD solutions (figure 4), the agreement is within experimental uncertainty for loading charge below 0.5, excepted at 318 K and w4MPD = 0.1687 where the calculated enthalpy of solution is 10% under estimated. At α > 0.5, the calculated enthalpies of solution are about 10% lower than experimental values, as observed for the 3MPD. As reported by Arcis et al.,24, 26 the accuracy on the calculated enthalpy depends mainly on the accuracy on the equilibrium constants of the reactions involved in gas dissolution. In this study the equilibrium constants of carbamate formation were not available for 3MPD and 4MPD; thus these values were adjusted in the model. These equilibrium constants for carbamate formation (eq 6), determined by the thermodynamic model, provide an accurate description of the relevant energetic properties of CO2 dissolution. Moreover the ability of the model to correlate enthalpy data can also be slightly influenced by the estimation of the activity coefficient. The approximations made on the Pitzer interaction parameters may explain the higher enthalpy deviations observed when increasing loading charges; non ideality increasing during gas dissolution because of ion species formation. As shown in figures 3 and 4, the enthalpy curve shape is mostly driven by the amine protonation (eq 5) and carbamate formation (eq 6) reactions, contributing at about 70% and



30% to the total energetic effect, respectively. The enthalpy of solution remains constant at low loadings up to α = 0.5 and then decreases. This trend is tightly linked to the mechanism of CO2 dissolution in secondary (or primary) amine.24, 25 In order to further understand the reaction mechanism, the extents of all the reactions involved during gas dissolution are highlighted in Figures 5 and 6 where the compositions of products of reactions between CO2 and amine are represented as function of gas loadings. For gas loading charge up to 0.5 the amines react to form protonated amines (eq 5) and carbamates (eq 6). Furthermore, the high pKa values of 3MPD and 4MPD cause a conversion of a fraction of the aqueous CO2 into carbonates (eqs 3 - 4), with a maximum at α = 0.5. Above this gas loading charge, equation 6 reverses because of the carbamate hydrolysis29 allowing the gas dissolution to continue with the formation of bicarbonates (eq 3, 4 and 6) and protonated amine species (eqs 5 and 6). These modifications in the mechanism of gas dissolution induce a slope change after α = 0.5 for the enthalpy of solution (Figure 1).

Influence of the position of the methyl group on the absorption of CO2 in aqueous solutions of methylpiperidine isomers. The thermodynamic model provides both a qualitative and quantitative representation of the {CO2 – methylpiperidine – H2O} system. It allows us to discuss the mechanisms involved in CO2 dissolution in aqueous solutions of isomers of methylpiperidine (2MPD, 3MPD and 4MPD). In a recent study,21 2MPD was observed to react as hindered amines i.e. to form unstable carbamate. Then the main contribution for the total enthalpy of solution was the reaction of amine protonation. In case of 3MPD and 4MPD, the contributions of carbamate formation reactions are significant. In Figure 7 are compared the average values for the enthalpies of solution and solubility limits of CO2 in 2MPD, 3MPD and 4MPD solutions. The properties are calculated for amine composition wa = 0.3, at 313.15 K and 0.11 MPa. These calculated enthalpies of solution in 2MPD, 3MPD and 4MPD are rather similar. However, a trend could be a slight decrease of the absolute value when the methyl group moves from the position 2 to position 4 on the piperidine ring. The CO2 solubility is higher in 2MPD than in 3MPD and 4MPD. The 2MPD being a hindered amine it reacts with CO2 as a tertiary amine. To further understand, the speciation curves derived from our model are given in Figure 8 and show how carbon dioxide is absorbed, with formation of carbamate, bicarbonate or carbonate. In 2MPD solution, carbon dioxide is absorbed as carbonates and bicarbonates. In 3MPD and 4MPD solutions


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CO2 is absorbed as carbamates and carbonate at loadings below 0.5. Increasing loading charge, carbamates are progressively replaced by bicarbonates. At 313.15 K, as mentioned previously, the formation of carbonates is found at low loadings (α < 0.5) because of the high pKa values of these isomers.

CONCLUSIONS Dissolution of CO2 in aqueous solutions of 3-methylpiperidine and 4-methylpiperidine, was investigated using a thermodynamic approach. Experimental enthalpies of solution and gas solubility data were determined as a function of temperature (318 K and 338 K), pressure (0.5 to 1.5 MPa) and amine composition. A rigorous model based on a γ–φ approach was developed to describe CO2 dissolution. Interaction parameters determined in a previous study to model {CO2 – 2MPD –H2O} system were used to represent the non-ideality of the liquid phase. The non-ideality of the vapor phase was described using the SRK equation of state. As there is no available data reported in the literature for the carbamate formation constant of either 3MPD or 4MPD, coefficients of a temperature dependent relation for these two equilibrium constants were determined in the regression procedure. Due to the complexity of the studied system some approximations were performed to estimate interaction parameters in solution. The modeling results are in good agreement with experimental solubility and enthalpy data. The thermodynamic representation makes it possible to analyze and compare the mechanisms and energetic effects of CO2 dissolution in aqueous solutions of different methylpiperidines (2MPD, 3MPD and 4MPD). It was shown that for 3MPD and 4MPD, the carbamate formation reaction contribute significantly to the total enthalpy of solution of CO2. This enthalpy of solution was found to be of the same magnitude as for 2MPD. Results from this modeling study contribute to the understanding of the underlying molecular mechanisms in CO2 dissolution in demixing solvents. The experimental enthalpies of solution (∆solH) and gas solubility (s) determined for these methylpiperidine derivatives are key parameters for the design of new phase-change solvents.



Acknowledgements: This work is realized with the financial support of ANR and NSERC through an international collaborative project between France and Canada named DACOOTA (n°ANR-12-IS09-000101).

ASSOCIATED CONTENT Supporting Information. In Part A are reported the tables of the chemicals (Table S1) and the experimental enthalpies of solution of CO2 (Tables S2-S8). Details on the thermodynamic modeling of CO2 dissolution in aqueous solutions of 3-methylpiperidine and 4methylpiperidine are given in Part B; all the model parameters are reported in Tables S9-S12.


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Chemical & Engineering Data 1993, 38 (3), 428-431. 10.1021/je00011a026 21.

Coulier, Y.; Lowe, A.; Tremaine, P. R.; Coxam, J. Y.; Ballerat-Busserolles, K.,

Absorption of CO2 in Aqueous Solutions of 2-Methylpiperidine: Heats of Solution and Modeling. International Journal of Greenhouse Gas Control 2016, 47, 332-329. https://doi.org/10.1016/j.ijggc.2016.02.009 22.

Arcis, H.; Ballerat-Busserolles, K.; Rodier, L.; Coxam, J.-Y., Enthalpy of Solution of

Carbon Dioxide in Aqueous Solutions of Monoethanolamine at Temperatures of 322.5 K and 372.9 K and Pressures up to 5 MPa. Journal of Chemical & Engineering Data 2011, 56 (8), 3351-3362. 10.1021/je2002946 23.

Span, R.; Wagner, W., A New Equation of State for Carbon Dioxide Covering the

Fluid Region from the Triple‐Point Temperature to 1100 K at Pressures up to 800 MPa.

Journal of Physical and Chemical Reference Data 1996, 25 (6), 1509-1596. doi:http://dx.doi.org/10.1063/1.555991 24.

Arcis, H.; Ballerat-Busserolles, K.; Rodier, L.; Coxam, J.-Y., Measurement and

Modeling of Enthalpy of Solution of Carbon Dioxide in Aqueous Solutions of Diethanolamine at Temperatures of (322.5 and 372.9) K and Pressures up to 3 MPa. Journal

of Chemical & Engineering Data 2012, 57 (3), 840-855. 10.1021/je201012e 25.

Arcis, H.; Coulier, Y.; Ballerat-Busserolles, K.; Rodier, L.; Coxam, J.-Y., Enthalpy of

Solution of CO2 in Aqueous Solutions of Primary Alkanolamines: A Comparative Study of Hindered and Nonhindered Amine-Based Solvents. Industrial & Engineering Chemistry

Research 2014, 53 (27), 10876-10885. 10.1021/ie501589j 26.

Arcis, H.; Rodier, L.; Ballerat-Busserolles, K.; Coxam, J.-Y., Modeling of (vapor +

liquid) equilibrium and enthalpy of solution of carbon dioxide (CO2) in aqueous methyldiethanolamine (MDEA) solutions. The Journal of Chemical Thermodynamics 2009,

41 (6), 783-789. http://dx.doi.org/10.1016/j.jct.2009.01.005 27.

Soave, G., Equilibrium constants from a modified Redlich-Kwong equation of state.

Chemical Engineering Science 1972, 27 (6), 1197-1203. http://dx.doi.org/10.1016/00092509(72)80096-4



28.

Edwards, T. J.; Maurer, G.; Newman, J.; Prausnitz, J. M., Vapor-liquid equilibria in

multicomponent aqueous solutions of volatile weak electrolytes. AIChE Journal 1978, 24 (6), 966-976. 10.1002/aic.690240605 29.

Sartori, G.; Savage, D. W., Sterically hindered amines for carbon dioxide removal

from gases. Industrial & Engineering Chemistry Fundamentals 1983, 22 (2), 239-249. 10.1021/i100010a016 30.

Kurz, F.; Rumpf, B.; Maurer, G., Vapor-liquid-solid equilibria in the system NH3-

CO2-H2O from around 310 to 470 K: New experimental data and modeling. Fluid Phase

Equilibria 1995, 104 (0), 261-275. http://dx.doi.org/10.1016/0378-3812(94)02653-I


Page 14 of 43

Page 15 of 43 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


LIST of TABLE

Table 1. Average values for the enthalpies of solution at low loadings (∆solHav) in aqueous 3MPD and 4MPD solutions at (318 and 338) K and their respective standard deviations (u). Table 2. Experimental values for the solubility (s) of CO2 in aqueous 3MPD and 4MPD solutions at (318 and 338) K and their respective uncertainties (u). Table 3. . Parameters used in Eq. (11) to represent the temperature dependence of the equilibrium constants of reactions 3-6.



Page 16 of 43

Table 1. Average values for the enthalpies of solution at low loadings (∆solHav) in aqueous 3MPD and 4MPD solutions at (318 and 338) Ka and their respective standard deviations (u). pb

-∆solHav

u(∆solHav)

kJ mol-1 of CO2

MPa

pb MPa

-∆solHav

u(∆solHav)

kJ mol-1 of CO2

T=318 K 0.43 0.96 1.47 0.43 0.96 1.48

0.46 0.99 1.51 0.47 0.96 1.47 a b

w3MPD = 0.1936 113.6 111.6 103.0 w4MPD = 0.1657 118.8 108.7 102.5 w3MPD = 0.4000 106.4 99.9 96.6 w4MPD = 0.4000 100.0 96.6 89.9

0.7 1.6 1.1

0.56 1.07 1.47

2.3 0.6 3.0

0.43 0.95 1.45 T=338 K

2.3 1.5 1.1

0.47 0.98 1.5

w3MPD = 0.4044 110.1 104.9 95.3 w4MPD = 0.3986 113.9 108.0 102.7 w4MPD = 0.2000 101.4 97.9 95.7

1.6 2.4 1.4

u(T)= 0.03 K u(p)= 0.02 MPa


1.1 1.6 1.6 1 1.2 1.8

1.3 1.4 0.1

Page 17 of 43 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Table 2. Experimental values for the solubility (s) of CO2 in aqueous 3MPD and 4MPD solutions at (318 and 338) Ka and their respective uncertainties (u). pb s u(s) pb s u(s) MPa

molCO2/molMPD

MPa

molCO2/molMPD

T=318 K 0.43 0.96 1.47 0.43 0.96 1.48

w3MPD = 0.1936 0.97 1.04 1.07 w4MPD = 0.1657 1 1.06 1.11

0.07 0.07 0.07

0.56 1.07 1.47

0.06 0.07 0.07

0.43 0.95 1.45

w3MPD = 0.4044 0.88 0.94 0.96 w4MPD = 0.3986 0.88 0.94 0.97

0.06 0.07 0.07 0.06 0.06 0.07

T=338 K

w3MPD = 0.4000 0.46 0.81 0.06 0.99 0.87 0.06 1.51 0.91 0.06

w4MPD = 0.2000 0.47 0.93 0.06 0.98 0.99 0.07 1.5 1.03 0.07

w4MPD = 0.4000 0.47 0.81 0.06 0.96 0.88 0.06 1.47 0.91 0.07 a b

u(T)= 0.03 K u(p)= 0.02 MPa



Page 18 of 43

Table 3. Parameters used in Eq. (11) to represent the temperature dependence of the equilibrium constants of reactions 3-6. Reaction number 3 4 5 5 6 6

Amine

A/K

B

C

Source

235.482 220.67 -2.2497 -2.6023 -9.8489 -7.5250

-36.7816 -35.5819 0 0 0 0

30

3MPD 4MPD 3MPD 4MPD

K-1 -12092.1 -12431.7 -6820.8 -6755.9 4344.5 3581.3


30 18

18

this work this work

Page 19 of 43 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


LIST of FIGURES

Figure 1: Enthalpy of solution (−∆solH) versus CO2 loading for aqueous solutions of amine at T = 318 K and p = 0.96 MPa: □, 3MPD (w3MPD = 0.1936) and ○, 4MPD (w4MPD = 0.1657). (a) ∆solH / (kJ·mol–1 of CO2), straight lines show the average values for the enthalpies of solution at low loadings (α < 0.4); (b) ∆solH / (kJ·mol–1 of amine).

Figure 2: Experimental solubility (s) of CO2 in aqueous solutions of amine vs total pressure. (a), 3MPD: , T = 318 K, w3MPD = 0.1936; , T = 318 K, w3MPD = 0.4044; , T = 338 K,

w3MPD = 0.4000. (b), 4MPD: , T = 318 K, w4MPD = 0.1657; , T = 318 K w4MPD = 0.3986; ,

T = 338 K, w4MPD = 0.2000; , T = 338K, w4MPD = 0.4000. Solid lines are the calculated

solubility from the model.

Figure 3: Enthalpies of solution (∆solH / kJ.mol-1·of CO2) versus CO2 loading for aqueous solutions of 3MPD: ○, experimental data.

From the thermodynamic model: (1), total

enthalpy of solution; enthalpic contribution from reaction: (2), amine protonation (eq. 5); (3), carbamate formation (eq. 6); (4), CO2 physical dissolution (eq. 9); (5), second ionization of CO2 (eq. 4); (6), first ionization of CO2 (eq. 3). (a): T = 318 K, w3MPD = 0.1936; (b): T = 318 K, w3MPD = 0.4044; (c): T = 338 K, w3MPD = 0.4000.

Figure 4: Enthalpies of solution (∆solH / kJ.mol-1·of CO2) versus CO2 loading for aqueous solutions of 4MPD: ○, experimental data.


enthalpy of solution; enthalpic contribution from reaction: (2), amine protonation (eq. 5); (3), carbamate formation (eq. 6); (4), CO2 physical dissolution (eq. 9); (5), second ionization of CO2 (eq. 4); (6), first ionization of CO2 (eq. 3). (a): T = 318 K, w4MPD = 0.1657; (b): T = 318 K, w4MPD = 0.4000; (c): T = 338 K, w3-MPD = 0.2000; (d): T = 338 K, w4-MPD = 0.4000.

Figure 5: Liquid phase composition (ni, number of moles) and pH (solid lines) calculated by the model versus gas loading charge of aqueous solutions of 3MPD. (a): T = 318 K and

w3MPD = 0.1936; (b): T = 318 K and w3MPD = 0.4044; (c): T = 338 K and w3MPD = 0.4000. Figure 6: Liquid phase composition (ni, number of moles) and pH (solid lines) calculated by the model versus gas loading charge of aqueous solutions of 4MPD. (a): T = 318 K and



w4MPD = 0.1657; (b): T = 318 K and w4MPD = 0.4000; (c): T = 338 K and w3-MPD = 0.2000; (d): T = 338 K and w4-MPD = 0.4000. Figure 7: Comparison at T =313.15 K of the calculated enthalpies of solution and solubility of CO2 in aqueous solutions of methylpiperidine isomers (wa = 0.3): 2MPD, 3MPD and 4MPD.

Figure 8: Fraction of CO2 absorbed in solution at equilibrium (T = 313.15 K, and wa = 0.3000) for aqueous solutions of methylpiperidine isomers derived from the thermodynamic model: (a), 2MPD 21; (b), 3MPD and (c), 4MPD.


Page 20 of 43

Page 21 of 43

120

(a) 110 100

-∆ ∆solH / kJ.mol-1

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


90 80 70 60 50 -2.22E-16 0.2

0.4

0.6

0.8

1

α (moleCO2/moleamine)


1.2

1.4


90

(b) 80 70


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 22 of 43

60 50 40 30 20 10 0 0

0.2

0.4

0.6

0.8

1

1.2

1.4

α (moleCO2/moleamine) Figure 1: Enthalpy of solution (−∆solH) versus CO2 loading for aqueous solutions of amine at T = 318 K and p = 0.96 MPa: □, 3MPD (w3MPD = 0.1936) and ○, 4MPD (w4MPD = 0.1657). (a) ∆solH / (kJ·mol–1 of CO2), straight lines show the average values for the enthalpies of solution at low loadings (α < 0.4); (b) ∆solH / (kJ·mol–1 of amine), straight lines show the slope change.


Page 23 of 43

2,500

(a) 2,000

1,500

p / kPa

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


1,000

500

0 0.5

0.6

0.7

0.8

0.9

1

α (moleCO2/mole3MPD)


1.1

1.2


2,500

(b) 2,000

1,500

p / kPa

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 24 of 43

1,000

500

0 0.5

0.6

0.7

0.8

0.9

1

1.1

1.2

α (moleCO2/mole4MPD) Figure 2: Experimental solubility limits (s) of CO2 in aqueous solutions of amine vs total pressure. (a), 3MPD: , T = 318 K, w3MPD = 0.1936; , T = 318 K, w3MPD = 0.4044; , T = 338 K, w3MPD = 0.4000. (b), 4MPD: (), T = 318 K, w4MPD = 0.1657; (), T = 318 K w4MPD = 0.3986; , T = 338 K, w4MPD = 0.2000; , T = 338K, w4MPD = 0.4000. Solid lines are the calculated solubility limits from the model.


Page 25 of 43

120

(a)

(1) 100


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


80

(2)

60 40

(3) 20

(4) (5) (6)

0 -20 0

0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9



1


120

(b)

(1)

100


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 26 of 43

80

(2) 60 40

(3)

20

(4)

0

(5) (6)

-20 0

0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9



1

Page 27 of 43

140

(c) 120

(1)

100


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


80

(2) 60 40

(3)

20

(4) (5) (6)

0 -20 0

0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9

1

α (moleCO2/mole3MPD) Figure 3: Enthalpies of solution (∆solH / kJ.mol-1·of CO2) versus CO2 loading for aqueous solutions of 3MPD: ○, experimental data.


enthalpy of solution; enthalpic contribution from reaction: (2), amine protonation (eq. 5); (3), carbamate formation (eq. 6); (4), CO2 physical dissolution (eq. 9); (5), second ionization of CO2 (eq. 4); (6), first ionization of CO2 (eq. 3). (a): T = 318 K, w3MPD = 0.1936; (b): T = 318 K, w3MPD = 0.4044; (c): T = 338 K, w3MPD = 0.4000.



140

(a) 120

(1) 100


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 28 of 43

(2)

80 60 40

(3) (4) (5) (6)

20 0 -20 0

0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9



1

Page 29 of 43

120

(b)

(1)

100


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


80

(2)

60 40

(3) (4)

20

(5) (6)

0 -20 0

0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9



1


120

(c)

(1) 100


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 30 of 43

80

(2)

60 40

(3) (4) (5) (6)

20 0 -20 0

0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9



1

Page 31 of 43

120

(d)

(1) 100


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


80

(2)

60 40

(3)

20

(4)

0

(5) (6)

-20 0

0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9

1

α (moleCO2/mole4MPD) Figure 4: Enthalpies of solution (∆solH / kJ.mol-1·of CO2) versus CO2 loading for aqueous solutions of 4MPD: ○, experimental data.


enthalpy of solution; enthalpic contribution from reaction: (2), amine protonation (eq. 5); (3), carbamate formation (eq. 6); (4), CO2 physical dissolution (eq. 9); (5), second ionization of CO2 (eq. 4); (6), first ionization of CO2 (eq. 3). (a): T = 318 K, w4MPD = 0.1657; (b): T = 318 K, w4MPD = 0.4000; (c): T = 338 K, w3-MPD = 0.2000; (d): T = 338 K, w4-MPD = 0.4000.



3.0

12

(a) RR'NH2+

2.5

11

10

RR'NH HCO3-

1.5

9

1.0

8

RR'NCOOCO32-

0.5

7

CO2

0.0

6 0

0.1

0.2

0.3

0.4

0.5

0.6

0.7



0.8

0.9

1

pH

2.0

ni / mol

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 32 of 43

Page 33 of 43

8.0

12

(b) 7.0

RR'NH2+

11

6.0

RR'NH

HCO3-

4.0

9

3.0

RR'NCOO-

8

CO32-

7

2.0 1.0

CO2

0.0

6 0

0.1

0.2

0.3

0.4

0.5

0.6

0.7



0.8

0.9

1

pH

10

5.0

ni / mol

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60



8.0

12

(c) 7.0

11 +

RR'NH2

6.0

RR'NH

HCO3-

4.0 3.0

9

RR'NCOO-

8

CO32-

7

pH

10

5.0

ni / mol

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 34 of 43

2.0 1.0

CO2

0.0

6 0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0.9

1

α (moleCO2/mole3MPD) Figure 5: Liquid phase composition (ni, number of moles) and pH (solid lines) calculated by the model versus gas loading charge of aqueous solutions of 3MPD. (a): T = 318 K and

w3MPD = 0.1936; (b): T = 318 K and w3MPD = 0.4044; (c): T = 338 K and w3MPD = 0.4000.


Page 35 of 43

2.5

12

(a) 11

RR'NH2+

2

HCO39

1 8

RR'NCOO0.5

7

CO32-

0 0

CO2

0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9



6 1

pH

10

RR'NH

1.5

ni / mol

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60



8.0

12

(b) 7.0

11

RR'NH2+

6.0

HCO3-

4.0

9

3.0

8

RR'NCOO-

2.0

CO321.0

7

CO2

0.0

6 0

0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9



1

pH

10

RR'NH

5.0

ni / mol

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 36 of 43

Page 37 of 43

3.0

12

(c) 2.5

11

+

RR'NH2

10

RR'NH

HCO3-

1.5

1.0

9

8

RR'NCOO-

0.5

7

CO32-

CO2

0.0

6 0

0.1

0.2

0.3

0.4

0.5

0.6

0.7



0.8

0.9

1

pH

2.0

ni / mol

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60



8.0

12

(d) 7.0

11

RR'NH2+

6.0

HCO3-

RR'NH

4.0

9

3.0

pH

10

5.0

ni / mol

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 38 of 43

8

RR'NCOO2.0

CO32-

1.0

7

CO2

0.0

6 0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0.9

1

α (moleCO2/mole4MPD) Figure 6: Liquid phase composition (ni, number of moles) and pH (solid lines) calculated by the model versus gas loading charge of aqueous solutions of 4MPD. (a): T = 318 K and

w4MPD = 0.1657; (b): T = 318 K and w4MPD = 0.4000; (c): T = 338 K and w3-MPD = 0.2000; (d): T = 338 K and w4-MPD = 0.4000.


1

110

s - ∆solH

0.95

s (moleCO2/molea)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


105 100

0.9

95

0.85

90 85

0.8

80

0.75


Page 39 of 43

75

0.7

70 2MPD

3MPD

4MPD

Figure 7: Comparison at T =313.15 K of the calculated enthalpies of solution and solubility limits of CO2 in aqueous solutions of methylpiperidine isomers (wa = 0.3): 2MPD, 3MPD and 4MPD.



1.00

(a)

2-

CO3

0.90

fraction of CO2 in solution

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 40 of 43

0.80

HCO3-

0.70 0.60 0.50 0.40 0.30 0.20 0.10

CO2

0.00 0

0.1

0.2

0.3

0.4

0.5

0.6

0.7



0.8

0.9

1

Page 41 of 43

1.00

(b)

0.90


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


3MPDCOO-

0.80

HCO3-

0.70 0.60 0.50 0.40

CO32-

0.30 0.20 0.10

CO2

0.00 0

0.1

0.2

0.3

0.4

0.5

0.6

0.7



0.8

0.9

1


1.00

(c)

0.90


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 42 of 43

0.80

4MPDCOO-

HCO3-

0.70 0.60 0.50 0.40

CO32-

0.30 0.20 0.10

CO2

0.00 0

0.1

0.2

0.3

0.4

0.5

0.6

0.7

0.8

0.9

1

α (moleCO2/mole4MPD) Figure 8: Fraction of CO2 absorbed in solution at equilibrium (T = 313.15 K, and wa = 0.3000) for aqueous solutions of methylpiperidine isomers derived from the thermodynamic model: (a), 2MPD 21; (b), 3MPD and (c), 4MPD.


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TOC/Abstarct Graphic. -

O

+

N

N

+

O

C

+

O

CH3

CH3

CH3

O

O

-

H

H

H N 2

H

H

H N 2

O

+

N

N

+

O

C

+

O

CH3

CH3

CH3

Synopsis. Carbon capture and storage is an important pathway towards the reduction of the impact of energy production on our environment.


Thermodynamic Modeling and Experimental Study of CO2 Dissolution

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